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02/24/10 Lecture
Announcements
• Corrections to Monday’s Powerpoint Lecture Slides Posted (Slides 3, 9 and 15)
• I have posted last semester’s Exam 1 plus the key
• Exam 1 – Topics– Ch. 1, 3, and 4 (all topics covered in
reading/lectures)– Ch. 6 (Sections 1 through 4 expected)
Today’s Lecture
• Chapter 6 Material– Thermodynamics
• Definitions• Determination of signs of ΔS, ΔH, ΔG• Relationship between ΔG, Q and K• Le Chatelier’s Principle
– Solubility Product Problems
Thermodynamics
Some Definitions:H = EnthalpyS = EntropyG = Gibbs Free EnergyT = Absolute temperature
Thermodynamics
1. ΔH is related to heat of reaction- if a reaction produces heat, ΔH < 0 and
reaction is “exothermic”- a reaction that requires heat has ΔH >
0 and is endothermic
2. ΔS is related to disorder of system- If the final system is “more random”
than initial system, ΔS > 0
Thermodynamics
Entropy Examples: (Is ΔS > or < 0?)H2O(l) ↔ H2O(g)
H2O(s) ↔ H2O(l)
NaCl(s) ↔ Na+ + Cl-
2H2(g) + O2(g) ↔ 2H2O(g)
N2(g) + O2(g) ↔ 2NO(g)
ΔS > 0
ΔS > 0ΔS > 0
ΔS < 0
ΔS > 0
Thermodynamics
ΔG = Change in Gibbs free energyThis tells us if a process is spontaneous
(expected to happen) or non-spontaneous
ΔG < 0 process is spontaneous (favored)ΔG = ΔH - TΔS (T is absolute temperature)
processes that are exothermic (Δ H < 0) and increase disorder (Δ S > 0) are favored at all Tprocesses that have Δ H > 0 and Δ S > 0 are favored at high T
Example question
The reaction N2(g) + O2 (g) ↔ 2NO(g) has a positive H. Under what conditions is this process spontaneous?- all temperatures- low temperatures- high temperatures- never
Thermodynamics
ΔG and EquilibriumΔG = ΔG° + RTlnQ Q = Reaction
Quotient(for A ↔ B, Q = [B]/[A])At equilibrium, ΔG = 0 and Q = K
ΔG° = -RTlnK
Thermodynamics
• Example Question:The ΔG° for the reactionCa2+ + 2OH- => Ca(OH)2(s) is -52 kJ/mol
Determine K at T = 20°C for Ca(OH)2(s) => Ca2+ + 2OH-
Section 6 – 2: Le Châtelier’s Principle
• The position of a chemical equilibrium always shifts in a direction that tends to relieve the effect of an applied stress
• Types of stresses:– Addition/removal of reactant/product– Change in volume (gases) or dilution (aqueous solutions)– Changes in temperatures
• Can take intuitive or mathematical approaches to solving problems
Le Châtelier’s Principle
Intuitive Method– Addition to one side
results in switch to other side
– Example:
Mathematical Method
AgCl(s) ↔ Ag+ + Cl-
Additional Ag+ results in more AgCl
QRTGG ln
K
QRTG ln
When Q>K, ΔG>0 (toward reactants)
When Q<K, ΔG<0 (toward products)
Example: Q = [Ag+][Cl-]
As Ag+ increases, Q>K
Le Châtelier’s Principle
Stress Number 1 Reactant/Products:
Addition of reactant: shifts toward product
Removal of reactant: shifts toward reactant
Addition of product: shifts toward reactant
Removal of product: shifts toward product
Le Châtelier’s Principle
Stress Number 1 Example:CaCO3(s) + 2HC2H3O2(aq) ↔ Ca(C2H3O2)2(aq) + H2O(l) + CO2(g)
1. Add HC2H3O2(aq)
2. Remove CO2(g)
3. Add Ca(C2H3O2)2(aq)
4. Add CaCO3(s) No effect because (s)
Le Châtelier’s Principle
Stess Number Two: Dilution
Side with more moles is favored at lower concentrations
Example: HNO2(aq) ↔ H+ + NO2-
If solution is diluted, reaction goes to products
If diluted to 2X the volume:
2
2
HNONOH
K
2
2
21
21
21
HNO
NOHQ
KQ2
1
So Q<K, products favored
Le Châtelier’s PrincipleStess Number Two: Dilution – Molecular Scale View
H+ NO2-
Concentrated Solution
H+ NO2-
H+
NO2-
H+ NO2-
H+ NO2-
Diluted Solution – dissociation allows ions to fill more space
H+ NO2-
H+ NO2-
H+
NO2-
H+ NO2-
H+ NO2-
Le Châtelier’s Principle
Stress Number 3: TemperatureIf ΔH>0, as T increases, products favoredIf ΔH<0, as T increases, reactants favoredEasiest to remember by considering heat
a reactant or productExample:
OH- + H+ ↔ H2O(l) + heatIncrease in T
Solubility Product Problems
Importance:- gravimetric analysis
- chemical separations (e.g. selective removal of Mg2+ or Ca2+ to determine single ion in water hardness titration)
Solubility Product Problems - Solubility in Water
Example: solubility of Mg(OH)2 in waterSolubility defined as mol Mg(OH)2 dissolved/L
sol’n or g Mg(OH)2 dissolved/L sol’n or other units
Use ICE approach:Mg(OH)2(s) ↔ Mg2+ + 2OH-
Initial 0 0Change +x +2xEquilibrium x 2x
Solubility Product Problems
- Solubility of Mg(OH)2 in water
Equilibrium Equation: Ksp = [Mg2+][OH-]2
Ksp = 7.1 x 10-12 = x(2x)2 = 4x3 (see Appendix F for Ksp)
x = (7.1 x 10-12/4)1/3 = 1.2 x 10-4 M
Solubility = moles Mg(OH)2 dissolved = x
Solubility = 1.2 x 10-4 MConc. [OH-] = 2x = 2.4 x 10-4 M
Solubility Product Problems
- Solubility of Mg(OH)2 in Common Ion
If we dissolve Mg(OH)2 in a common ion (OH- or Mg2+), from Le Châtelier’s principle, we know the solubility will be reduced
Example 1) What is the solubility of Mg(OH)2 in a pH = 11.0 buffer?
Ksp = 7.1 x 10-12
Solubility Product Problems
- Solubility of Mg(OH)2 in Common Ion
Example 2) Solubility of Mg(OH)2 in 5.0 x 10-3 M MgCl2.
Solubility Product Problems Precipitation Problems
What occurs if we mix 50 mL of 0.020 M BaCl2 with 50 mL of 3.0 x 10-4 M (NH4)2SO4?
Does any solid form from the mixing of ions?
What are the concentrations of ions remaining?
Precipitations Used for Separations
Example: If we wanted to know the concentrations of Ca2+ and Mg2+ in a water sample. EDTA titration gives [Ca2+] + [Mg2+]. However, if we could selectively remove Ca2+ or Mg2+ (e.g. through titration) and retitrate, we could determine the concentrations of each ion.
Determine if it is possible to remove 99% of Mg2+ through precipitation as Mg(OH)2 without precipitating out any Ca(OH)2 if a tap water solution initially has 1.0 x 10-3 M Mg2+ and 1.0 x 10-3 M Ca2+.